laboratory practices in progress

7/29/2019 LABORATORY Practices in Progress

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NATIONAL UNIVERSITY OF RWANDAFACULTY OF MEDICINE

DEPARTMENT OF PHARMACY

L ABORATORY PRACTI CES OFCHEMISTRY

1. SAFETY AND L ABORATORY RULES2. COMM ON L ABORATORY MATERIAL S AND THEIR USAGES3. PREPARATI ON OF SOLUTIONS4. TI TRATIONS5. ORGANIC CHEMI STRY

K AGISHA Vdaste, MSc

Academic year 2010
http://www.go2pdf.com/


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Laboratory practices: Pharmacy department (Academic year 2010)

Chapter I: SAFETY AND L ABORATORY R ULES.

I .1Attitudes and Preparation

a. Come to laboratory periods on time and mentally prepared by studying the experiment andplanning your activities. This will help us to interact with each other.b. Be prepared physically; for example, don't try to do lab work on an empty stomach, or withoutsleep or when you have alcohol in your bloodstream.c. Write everything you do and see in your notebook so that you can trace your actions and makecorrections if necessary.d. Wear sensible clothing, including shoes that are comfortable and permit rapid movement incase of emergency, and hair or hat that does not obstruct your view or dangle into theexperiment.e. If you wear contact lenses, try to avoid wearing them in the lab. If you must wear contactlenses, your goggles must seal particularly well to your face.f. If you injure yourself, even slightly, report it to the safety-officer and/or your instructor, andseek first aid. If you experience eye irritation, flush your eyes at the nearest emergency eyewashstation for 15 minutes (remove contacts) and seek medical attention immediately.g. If you have any existing physical conditions that might affect your performance, your health,or other peoples' health in the lab, please inform your instructor. This information will be keptconfidential; examples might include pregnancy, medications, allergies, epilepsy, Specialarrangements may be possible.

I .2 Y our WorkingEnvironment

a GOGGLES for chemical splash protection are required at all times in labs or instrumentrooms, i.e. all parts of the lab, even when your are not handling chemicals. The goggles willprotect your eyes from most splashes and impacts. The goggles do not meet the standard if theair baffles are removed. Some people have trouble with their goggles fogging up. The bestsolution is to take a short break outside the lab to clean them.b. Rubber gloves are strongly recommended to protect you from absorption of chemicals throughthe skin. We also recommend a lab coat to protect your clothes and skin from your and yourneighbor's spills.c. Keeping your bench space tidy will minimize breakage and spills of your valuable products.

d. You are expected to clean up your own mess in community areas such as the IR room.e. Keep your glassware and other equipment cleaned up as you work. Having clean, dryglassware available at all times will save you much time in the long run.f. Be careful not to contaminate reagents with your spatulas or droppers. If you take too much ofa reagent, give it to a needy neighbor - do no return it to the bottle.g. Do not wander off with the only bottle of a reagent that everyone needs; keep it in its assignedlocation.h. Be sure the aisles are passable.


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I .3. Glassware

a. The most common laboratory injury is a cut occurring upon breakage of glass or porcelain.Most cuts can be prevented by careful work which prevents breakage.b. The safe procedure for inserting a glass tube or thermometer into a stopper with a hole is asfollows:i. Be sure the tip of the tube is fire-polished.ii. Lubricate the glass with glycerol or water.iii. Be sure the hole in the stopper is large enough.iv. Grasp the glass about 1" (no farther) from the end and push and twist to insertit into the stopper.v. Be sure that the hand holding the stopper is not in line with the entering glass.vi. As the glass begins to slide into the rubber, move the hand holding the glass back a little,always keeping it no more than 1" from the rubber.

vii. Most accidents occur because the glass snaps above the stopper from a force sideways(torque). Keeping your hand close to the stopper will help prevent your exerting a force sidewayson the glass.viii. The above considerations apply also to attachingrubber hoses to condensers.The condenser should be in your hand (not clamped to an apparatus) and gripped close to thelubricated connector being inserted into the hose.c. Never use a thermometer as a stirrer! Always support a thermometer in a beaker or flask witha clamp. If a mercury thermometer breaks, immediately contact the laboratory instructor andrestrict access to the area of contamination until cleanup can be arranged.d. Round-bottomed flasks will not stand upright by themselves and if rested on the counter willroll. They must be supported on a cork ring, in (not on) a beaker, or in a clamp.

e. When glassware is assembled, care should be taken to use the minimum number of clampsneeded for support, making sure:i. The clamp is attached to a vertical support bar.ii. No torque is applied by the clamp.iii. Top-heavy apparatus is prevented from rotating and tipping.iv. Hanging pieces are clamped - grease will not hold them against the force of gravity!f. Do not use a glass stopper to seal a hot container or you may never get it out again.Cork is recommended for organic solvents since rubber dissolves in organic solvents and vice-versa.g. Graduated cylinders are metastable and tip easily with the touch of a sleeve.h. Report breakage of glassware to your instructor for disposal instructions.

i. Think before cleaning equipment - it makes little sense to scrub a graduated cylinder thatcontained ether or a water-insoluble material with soap and water.

I .4Safety Equipment

a. Fire Extinguishers for smothering fires. Departments policy regarding response to firesrestricts the use of fire extinguishers to persons who are properly trained. Small fires may beextinguished by covering with a book or larger container.b. Fire Blanket for smothering fires.


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c. Safety Shower for rinsing chemicals off the body.d. Eye Wash Fountain for rinsing chemicals from the eyes.e. First Aid Kit - Note: even minor injuries must be reported to your instructor.f. At least two exits.g. Dustpan and broom for removing broken glass.

I .5Reagent

a. Assume that a particular reagent is hazardous unless you know for sure it is not.b. Never taste a chemical unless specifically directed to do so.c. If you are instructed to smell a chemical, point the vessel away from your face and carefullyfan the vapors toward your face with your hand and sniff gently.d. Dilute concentrated acids and bases by pouring the reagent into water (room temperature orlower) while stirring constantly. Never pour water into concentrated acids; the heat of solution

will cause the water to boil and the acid to splatter.e. Avoid rubbing your eyes unless you know your hands are cleanUse the fume hoods. Any experiment involving the use of or production of poisonous orirritating gases must be performed in a hood.f. Read the label. Read the label carefully, read it twice, before taking anything from a bottle.Many chemicals have similar names, such as sodium sulfate and sodium sulfite. Using the wrongreagent can spoil an experiment or can cause a serious accidentg. Be aware of your lab neighbors' activities; you may be a victim of their mistakes. If youobserve improper techniques or unsafe practices:h. Advise your neighbor. And advise your instructor if necessary.

I .6T oxic Hazards

a. The materials used in the organic lab are the safest we can find consistent with your need todevelop skills in working with hazardous materials in your career in science.b. Since you are wearing eye protection, the opportunity for liquids or solids to enter the eye issmall. Chemicals in the eye should be immediately flushed with copious amounts of water usingthe eyewash fountain.c. To prevent inhalation of organic and inorganic vapors, do your experiments in the fume hoodor under the minihoods on the bench.d. If your need to determine the odor of any material, waft it gently toward your nose with yourhand - don't stick your nose in the container and inhale.

e. Organic compounds can be absorbed through the skin, so be careful about spilling things.Wear rubber gloves to prevent contact with your skin, but treat the gloves as if they were bareskin, keeping them scrupulously clean. You might set aside a pen for laboratory work tominimize the possibility of contamination from your gloves via your pen to your hands and face.Obviously, chewing a pen or pencil that has been used in the lab would unwise.f. Organic vapors also can be absorbed into food or tobacco which you may ingest later.Moreover, any drinks brought into the lab could have things spilled into them. No food or drinksin the laboratory, not even stuffed in your backpack. If you do not have a locker to keep food in,please remove the food, drink and cigarettes to the hallway, or ask your instructor for a safeplace to keep it. Smoking is not allowed in


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State buildings, as the nicotine and other contents of the smoke are well-known health hazards(look up the LD50 of nicotine if you are skeptical).g. If you spill a liquid on the bench, immediately soak it up with paper towels and, if itis volatile, transfer the towels to the hood. Inform your instructor as to the nature of the spill incase further action is warranted.h. If concentrated acid is spilled, add sodium carbonate or bicarbonate, solution or solid.If concentrated base is spilled, add dilute and/or weak acid (e.g. acetic). Indicator solution orpaper will be available in the lab. If your skin (or clothing) comes in contact with the spill,immediately flush the skin or clothing with water for 15 minutes.i. Should you spill bromine solution anywhere, treat the spill immediately with sodiumthiosulfate solution.j. Bottles of the reagents mentioned in g) and h) are available on the small counter above yourbench.

I .7Heat Hazards

a. Most organic compounds are flammable and may catch fire even in the absence of flame athigh temperatures.b. Flames are rarely allowed in the organic laboratory. If flames are permitted by your instructor,plan your experiments so that you never leaveyour flameunattended.c. If you light a flame, you are responsible for the consequences, so check with your instructorfor a safe location.d. If you use a bunsen burner, be sure to tie back your hair and be careful that hair or clothing arekept clear of the flame.e. If there is a flame in the neighborhood, do not pour flammables; organic vapors are usually

denser than air and will flow along the bench without alerting you by their odors.f. Make sure you know the location of the nearest fire extinguisher and the nearest exit.g. Reactions that are exothermic or are being heated must be monitored; do not leave themwithout having someone watch.h. Never, never, never heat a closed system! Pressure will build up and cause the glass to fail,sending projectiles of glass in all directions. Do not depend on small leaks a substantial air exitmust be provided.

I.8If Thereisa Fire

a. In the lab where you are working.

i. Shout "fire" to alert your neighbors and instructor if you discover it.ii. A small fire in a test tube or other container can usually be extinguished by covering thecontainer with a watch glass or book. If the fire cannot be extinguished by one extinguisher or bysand or water, you will be instructed to evacuate, following the procedure in b).iii. One terrible possibility is that someone's clothing is set on fire. If the person runs, the flamewill be increased by increasing the supply of oxygen. It must be smothered. Wrap the person in alab coat, fire blanket, or whatever is handy to exclude oxygen.b. Elsewhere in the building (fire alarm sounds):i. Extinguish any flames and turn off electrical equipment.ii. Close any open windows and internal doors near you.


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iii. Walk quickly through the nearest exit to the hallway and leave the building by the neareststairwell.iv. The last person leaving the room, usually your instructor, will close the hall door.v Know the ways to put out a fire.

a) If it is open fire, such as a large chemical spill on a lab bench, the correct extinguisher shouldbe used as follows:

Pull the pin. Point the extinguisher (of dry) or hose (if CO2) at the base of the fire. Squeeze the handle while moving the extinguisher back and forth.

NOTE: Be careful not to spread the fire by getting the nozzle of the extinguisher too close--thematerial beingemitted is under pressure.

I .9L aboratory Electrical Equipment

a. During your s studies you will use a variety of instruments to analyze your samples.As with all electrical equipment, a certain amount of care is needed to prevent fire, shock anddamage to the equipment. Be careful not to bring water, especially on your hands, into contactwith connected electrical equipment.b. The hot plates you are provided are powerful and seldom need to be set higher thanc. Much of the heating in organic chemistry is done with electrical heating mantles; these mustbe plugged into a variable transformer, not directly into the outlet or they will overheat and maycause a fire.d. Never transfer anything into a flask that is sitting in a heating mantle; use a cork ring, beakeror clamp to hold the flask during transfers. Organics spilled in a mantle will catch on fire whenthe electricity is connected, acids or bases will corrode the wires, and water will cause a short

circuit.e. Never pour into a container on an electronic balance - they often have the wiring and knifeedges under the pan are thus easily damaged.f. Turn off electrical equipment immediately after you have finished unless your instructor hasstated otherwise (e.g. the gas chromatographs must be left on for an hour to stabilize).g. Report frayed cords, or non-functional equipment to your instructor. Do notput it back in thecupboard or you will be stuck with it again next time.h. No samples are allowed on top of any instruments.

I .10 Pressure Hazards

a. Never heat a closed system.b. When using a separatory funnel, vent frequently and remove the stopper immediately uponsetting it upright for separation.c. Compressed gas cylinders must be strapped to the bench above their center of gravity whenthe protective caps are off. Pressure regulators are generally not interchangeable between gasesfor safety reasons. Gas cylinders should be free of regulators and protected by their cap beforemoving.


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I .11 WasteDisposal

a. In order to minimize damage to the environment, and in compliance with Statechemical wastes must be separated into categories and carefully labeled as to their contents.Please read and follow the labels on the waste bottles to ensure that your chemical wastes aretreated safely and appropriately. You will find containers for:i. General Organic Waste (flammable)ii. Halogenated Hydrocarbons (non-flammable)iii. Chromic Acid Solutions (these have been phased out)iv. Leadv. Silvervi. Other Heavy Metalsvii. Waste from specific experiments in some cases.viii. Acids

ix. Basesx. In some experiments, acids and bases will be neutralized to a pH of 6 - 10(State law) as part of the experiment and flushed down the drain with lots of water. Yourinstructor will give you instructions in particular cases. Indicator solution or paper will beavailable in the lab.xi. Broken thermometers create the special problem of spilled mercury (a toxic heavy metal).Report such accidents immediately to your instructor; usually any mercury which cannot becollected is reacted with sulfur or absorbed with a special kit before disposal as heavy metalwaste.xii. Broken glass or porcelain is swept up into a dust pan and disposed of in a special containerfor broken glass. Please don't use your fingers.

I.12Summary

1.Approvedsafetygogglesmust bewornatall times.2.Nofood,drinksor smokingareallowed.3.Shoesmust beworn.Nobarefeet or thongsandalsareallowed.4.Nolab-work ispermittedwhenyouarealoneinthelab,speciallywithanysolventsor chemicals.5.Noopenflamesareallowedexcept asdirectedbytheinstructor.

6.Informyourself aboutthelocationof fireextinguishers,safetyequipment,emergencyshowersandeye-wash,andtheemergencytelephonenumbersandthenearest exit.7.Nounauthorizedexperimentsmaybeperformed.8.Donot usebrokenor crackedglassware.Checkglasswarebeforeusingit.9.Never tasteor smell chemicals.Consult chemical-safetyhandbookbeforeusinganyunknownchemical(whichistobeusedfor thefirst time).10.Avoidcontact ofchemicalswithskin.Theuseofrubber glovesisrecommended.havelabcoatesonyouall thetime11.Disposeofchemical wasteasdirectedbychemical-safetyhandbookor instructor.12.Cleanyour workareaandput awayall equipmentandglasswarebeforeleaving.makesureequipment isputawayin thecorrect locker- your personal locker or thecommonlocker.

13.Put paper trashandbrokenglassin trashcontainers.14.Keepinstrumentroomcleanandfreeofpaper.


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CHAPTER II : COMM ON L ABORATORY MATERIAL S AND THEIR USAGE

Distilling flasks: The figure (a) is the ordinary distilling flask.The sizes vary between 25 and5000ml.(b) is the so-called Claisen flask, a distilling flask with two necks; the thermometer is placed inthe neck currying the side arm. Sizes vary between 25 and 2000ml. It is of particular value indistillation where foaming or bumping occurs and is widely employed in distillations underdiminished pressure.(c) is identical with (b) except that it is provided with a second long and indented neck. It issometimes termed a Claisen flask with fractionating side arm.In(d) the side arm outlet extends a short distance into the long neck of the flask , thus preventingany vapour which has been in contact with cork or rubber stoppers from condensing and flowingdown sides the arm.

It is usually employed for those liquids which attack cork or rubbers stoppers.

2.1CONDENSERThe various types in common use are shown in the figure bellow.Type(a) is a typically L iebig condenser,which consists of an inner glass tube surrounded by aglass jacket trough which water is circulated. The inner jacket is fitted into the outer jacket bymeans of rubber stoppers; rubber tubing used formerly, is less durable and is not recommended.(b) is an all-glass Liebig condenser of similar design to (a); the jacket is sealed to the condensertube . Two convenient sizes of suitable for general use have jacket of 20 and 40cm length.In the Pyrex glass West condenser greater efficiency of cooling is obtained by having a light-

walled inner tube and a heavy-walled outer tube with a minimum space between them.(c) is the inner tube ofL iebig condenser. It is used as an air.condenser when the boiling point ofthe liquid is above 140-1500

(d) is an all-glass Allihn condenser. The condensing tube is made with a series of bulbs; thisincreases the condensing surface and lessens the resistance to the passage of vapours when thecondenser is employed for refluxing.(e) is a typical double surface condenser( Davies type). It is far efficient more than any of thepreceding types and the jacket is usually shorter


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(f) is an efficient spiral condenser of Friedrich type. The hot vapours can be introduced either atthe side or the bottom, thus permitting the use of the condenser either for condensing vapoursfrom another reaction vessel or for ordinary reflux purposes.(g) is coil condenser provided with an internal glasss coil or spiral. Ina modification there isboth an internal spiral as well as an outer cooling jacket.(h) is a Dewar type of refluxing condenser. It is usually charged with a freezing mixture, e.g.,Dry ice mixed with alcohol or acetone.


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2.2 VARIOUS K IND OF FUNNEL

Type (a) an ordinary filtration funnel having a 600 angleType (b) a wide-stemmed funnel is used when transferring powders.Type (c), (d)and (e) are known as reparatory funnels;

2


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.2.1FUNNEL SUITABLE FOR FILTRATION

The Buchner funnel (a) is used in conjunction with a filter flask into which is fitted by means of

rubber stopper. The use of either a flat annular rubber ring or a cone to provide a seal betweenflask and funnel as shown in (c) or (d) respectively is often more convenient.The Hirsch funnel (b) has sloping sides and is designed to deal with a smaller amount ofprecipitate than the Buchner funnel.The funnel in the assembly (d) is a substitute for the Hirsh funnel. It consists of an ordinary glassfunnel fitted with a witt plate(e) which is a perforated porcelain plate 1-4 cm of diameter, uponthe filter paper can rest.The slit sieve funnel (f) is constructed entirely of glass (Jena or Pyrex) and therefore possesobvious advantages over opaque (porcelain) Buchner or Hirsch funnel. Similar advantages areapparent with the sintered glass funnel (g), which is available in a number of porosities coarse,medium and fine)

2.3 APPARATUS FOR MEASURING MASS

Mass is measured using weighing balances. There are different types of weighing balances suchas beam balances and electrical balances


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2.4 APPARATUS FOR MEASURING TEM PERATURETemperature is measured using thermometers. There are different types of thermometers. Thefigure below shows the common type used in chemistry laboratories2.5 APPARATUS FOR MEASURING TI METhe apparatus for measuring time are usually watches and clocks. For accuracy duringexperiments in the laboratory, stop watches and stop clocks are used. Some of the common typesare shown in the figure bellow


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2.6 OTHER COMMONLY USED APPARATUS ARE REPRESENTED BELOWDIAGRAMMATICALL Y IN THE FIGURE:


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Name of Apparatus Use

1 Separating Funnel Seprating immiscible liquids

2 Round bottomed flasks Used when heating liquid substances

because heat is suplied uniformly

3 Flat-bottomed flask Used for general laboratory experiences

4 Gas jar Used for gas collection

5 Deflagratiing Used for holding burning substances

6 Trough Used foe holding somer amount of water for

some experiments

7 Thistle Used for delivering liquid substances

8 Desicccator Used for drying sustances or keeping

substances free from moisture

9 Boi ling -tube Used when heating liquid or solid

substances

10 Test-tube Used for general laboratory experiments

11 Test-tube rack Used for holding boiling tubes and test-tubes

12 Teat pipette(dropper) Used for delivering liquid substances drop-wise

13 Funnel Used for deli vering liquids carefully into vessels

14 Wash battle Used to hold water for rinsing of vessels

15 conical flask Used for general laboratory experiments

16 test-Tube holder Used for holding hot test-tubes

17 Pipe-clay triangle Used for supporting crucibles during heating

18 Evaporiting dish Used when evaporating liquids

19 Spatula Used for scooping solid substances from containers

20 Tongs Used for holding metal lic or non-meatllic substances

Used for storing chemicals

21 Wire gauze Used when glas is being heated or burned.

Causes even distribution of heat when heating

substances in beakers or flasks

22 Reagents bottles Used for storing chemicals

23 Clamp and stand Used for supporting and holding pieces of apparatus

during experiments

24 Crucible Used when heating solid substances that require strong heating

25 Tripod stand Used for supporting beakers and flaski in which liquids

are being heated

26 Pestle and mortar Pestle is used for crushing substances while the mortar

holds the substances being crushed


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CHAPTER II I: PREPARATION OF SOL UTIONS

3.1BASIC CONCEPTS3.1.1 EXPRESSIONS OF CONCETRATI ON

In chemistry, a solution is a homogeneous mixture composed of two or more substances. In such

a mixture, a solute is dissolved in another substance, known as a solvent. An aqueous solution

is a solution in which the solvent is water.

Concentration is the measure of how much of a given substance (solute) there is mixed with

another substance (solvent water). This can apply to any sort of chemical mixture, but most

frequently the concept is limited to homogeneous solutions, where it refers to the amount of

solute in a substance.

For scientific or technical applications, concentration is expressed in a quantitative way.

There are a number of different ways to quantitatively express concentration; the most common

are listed below. They are based on mass orvolume orboth. Depending on what they are based

on it is not always trivial to convert one measure to the other, because knowledge of the density

might be needed to do so. At times this information may not be available, particularly if the

temperature varies.

Mass can be determined at a precision of ~ 0.1 mg on a routine basis with an analytical balance.

Both solids and liquidsare easily quantified by weighing.

Some units of concentration -particularly the most popular one (molarity)- require to express the

mass of substance in the moles. The mole (symbol: mol) is the SI basic unit that measures an

amount of substance.

A mole is the amount of substance of a system which contains as many elementary entities as

there are atoms in 0.012 kilogram (or 12 g) of carbon-12 (12C), where the carbon-

12 atoms are unbound, at rest and in their ground state. The number of atoms in 0.012 kg of12C is known as the Avogadro constant (NA), and is determined empirically. The currently

accepted value is NA = 6.02214179(30)1023 mol-1.

When the mole is used to specify the amount of a substance, the kind of elementary entities

(particles) in the substance must be identified. The particles can be atoms, molecules, ions, etc.

The atomic mass of a chemical element is the mass of NA atoms (=1 mol) of it. The atomic

masses of the elements are included in the periodic table of the elements.


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The molar massof a molecule is the mass ofNA such molecules. The molar mass of a

molecule is equal to the sum of the atomic masses of its constituting atoms.

For example, table salt is sodium chloride (NaCl). The atomic mass of sodium, given in the

periodic table is 22.990 g/mol, and of chlorine is 35.453 g/mol. The molar mass of NaCl is,

therefore:

M(NaCl) = 22.990 + 35.453 = 58.443 g/mol.

The molar mass M of a molecule, multiplied by the number of moles n, is equal to the total mass

m(g) of the molecules:

m= nM [1]

In contrast to mass, a substance's volume is variable depending on ambient temperature andpressure. The volume of a liquid is usually determined by calibrated glassware such as burettes

and volumetric flasks. For very small volumes precision syringes are available.

Volumetric flaskVolumetric flasks (Fig.1; V=1l, 250 ml, 100ml, 5 ml, etc) are calibrated at a standard state

temperature and pressure (25oC, 101.325 kPa). (The measurement of mass does not require such

restrictions.) The use of graduated beakers and cylinders is not recommended as their indication

of volume is mostly for decorative rather than quantitative purposes. The volume of solids,

particularly of powders, is often difficult to measure, which is why mass is the more usual

measure.


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3.1.2CONCENTRATION:

Molar concentration (Molarity) denotes the number of moles of a given substance per litre of

solution.

Following the SI system of units, the National Institute of Standards and Technology, the

United States authority on measurement, considers the term molarity and the unit symbol M to

be obsolete, and suggests instead the 'amount-of-substance concentration' (c) with units mol/m3

or other units used alongside the SI such as mol/l. This recommendation has not been generallyimplemented in academia yet.

When discussing the molarity of minute concentrations, such as in much pharmacological

research, molarity is sometimes expressed in milimolars (mmol/l) or micromolars (1 millionth of

a molar).

3.2TO MAK E SOLUTION OF GIVEN MOLARITY AND VOL UME

1. Find the FW of the solute, usually from label.

2. Determine the molarity desired.

3. Determine the volume desired.

4. Determine how much solute is necessary by using the formula.

5. Weigh out the amount of solute.

6. Dissolve the solute in less than the desired final volume of solvent.

Place the solution in a volumetric flask or graduated cylinder. Add solvent until exactly the

required volume is reached, BringTo Volume, BTV

EXAMPLEHow much solute is required to make 300 mL of 0.8 M CaCl2?ANSWERFORMULA

FW X molarity x volume =g soluteneeded

111.0g) (0.8mole) (0.3L) =26.64 g

V

nC =


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mole L

Molal concentration, (Molality) denotes the number of moles of a given substance (n) per

kilogram ofsolvent (ms)(not solution). The SI unit for molality is mol/kg.

The determination of molality only requires a good balance, because the masses of both solvent

and solute can be obtained by weighing. Using a balance is often more precise than working with

volumetric flasks burettes and pipettes. Another advantage of molality is that it does not change

with the temperature as it deals with the mass of solvent, rather than the volume of solution.

Volume typically increases with increase in temperature resulting in decrease in molarity.Molality of a solution is always constant irrespective of the physical conditions like temperature

and pressure.

In a dilute aqueous solution near room temperature and standard atmospheric pressure, the

molarity and molality will be very similar in value. This is because 1000 g of water roughly

corresponds to a volume of 1 l at these conditions (supposing the density of water


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1 g/ml), and because the solution is dilute, the addition of the solute makes a negligible impact

on the volume of the solution. However, in all other conditions, this is usually not the case.

Themolefraction ?i, (also called molar fraction) denotes the number of moles of solute as a

proportion of the total number of moles in a solution.

Mole fractions are dimensionless quantities. For instance: 2 mole of solute A is dissolved in 6

moles of solvent B. Mole fractions:

This measure is used very frequently in the construction of phase diagrams. It has a number ofadvantages: the measure is not temperature dependent, does not require knowledge of the

densities of the phase(s) involved, a mixture of known mole fraction can be prepared by

weighing off the appropriate masses of the constituents. As both mole fractions and molality are

only based on the masses of the components it is easy to convert between these measures.

Other, common used expressions of concentration are:v Mass percentage (denotes the mass of a substance in a mixture as a percentage of the

mass of the entire mixture);

v Mass -volume percentage,

v Volume-volume percentage.

3.3 WEIGHT/VOLUME %Grams of solute100 mL total solution

EXAMPLE120 g of NaCl in

100 mL of total solution= 20% (w/v) solution.

EXAMPLE2How would you prepare 500 mL of a 5 % (w/v) solution of NaCl?ANSWERBy equation

1. Total volume required is 500 mL

2. 5% = 0.05

=

n

i

ii

ni

nX

25.062

2=

+=

+=

BA

AA

nn

nX


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3. (0.05) (500 mL) = 25

4. 25 is the amount of solute required in grams.

5. Weigh out 25 g of NaCl. Dissolve it in less than 500 mL of water.

6. In a graduated cylinder or volumetric flask, bring the solution to 500 mL.

TWO OTHER FORM S OF %v/v mL solute

100 mL solution

w/w g solute

100 g solution

3.4WEIGHT/WEIGHT

How would you make 500 g of a 5% solution of NaCl by weight (w/w)?ANSWER

1. Percent strength is 5% w/w, total weight desired is 500g.

2. 5% = 5g/100g

3. 5g X 500 g = 25 g = NaCl needed

100 g

4. 500 g 25 g = 475 g = amount of solvent needed


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Dissolve 25 g of NaCl in 475 g of water

PARTS

Parts may have any units but must be the same for all components of the mixture.EXAMPLE:A solution is 3:2:1 ethylene:chloroform:isoamyl alcoholMight combine:

3 liters ethylene

2 liters chloroform

1 liter isoamyl alcohol

PPM AND PPBv ppm: Thenumber of parts of soluteper 1 million parts of total solution.

v ppb: Thenumber of parts of soluteper billion parts of solution.

PPM EXAMPLE:5 ppm chlorine = 5 g of chlorine in 1 million g of solution, or 5 mg chlorine in 1 million mg of

solution, or 5 pounds of chlorine in

1 million pounds of solution

3.5 PPM TO MI CROGRAMS/mL

For any solute:1 ppm in water = 1 microgram

mL

Each star represents 1 mg of dioxin.

What is the concentration of dioxin in tube expressed as ppm (parts per million)? ____________

What is the total amount of dioxin in beaker? ___________


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Each star represents 1 mg of dioxin.

What is the total amount of dioxin in tube? 25 mg

What is the concentration of dioxin in tube expressed as ppm? ____________

1 ppm in water = 1 g

mL

25 mg/500 mL = 0.05 mg/mL = 50 g/mL

So the concentration is 50 ppm

3.6M AKI NG DIL UTIONS

In common chemical practice, one frequently needs to prepare a new solution by dilution of the

stock solution. This can be done following the mixing rule

C0V0 ??C1V1

where V0 is a volume of the stock solution at concentration C0, and V1 is a required volume of

a final solution with its final concentration C1.

The following example shows how to prepare 50 ml (V1) of 0.1 mol/l (C1) solution by dilution of

the stock solution at concentration C0=1 mol/l: According to the mixing rule, we need the

volumeV0 of a stock solution with concentration C0

C0V0 =?OC1V1

0

110

C

VCV =

mllmol

lmomlxV 5

/1

/1.0500 ==


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Therefore, we pour a small amount of the stock solution into the beaker and transfer 5 ml of it

into 50 ml volumetric flask using the pipet. 0.1 mol/l solution is then obtained by filling the flask

with distilled water up to the 50 ml mark.

3.6STOCK SOLUTION

We define a stock solution as a concentrate, that is, a solution to be diluted to some lowerconcentration for actual use. We may use just the stock solution or use it as a component in amore complex solution. We refer to the solution that we end up using as a working solution. Ifyou are comfortable making dilutions then you can appreciate the many advantages of workingwith stock solutions. Although it is never absolutely necessary to use a stock solution, it is oftenimpractical not to use them. Stock solutions can save a lot of time, conserve materials, reduce

needed storage space, and improve the accuracy with which we prepare solutions and reagents.Here are several illustrated types of applications using stock solutions.

How do I make a 10 millimolar solution of HCl ?It is easily you only need concentrate HCl and then dilute it into certain volume for desired

molarity. Concentrate HCl with percentage of 38% HCl has molarity about 12.39 M, assume that

you will make 10 millimolar (mM) HCl solution with 1 L in volume thus how many mL of

concentrate HCl do you need?

Just simple calculate using the dilution formula:

V1xM1 = V2xM2

10 mM = 0.01 M thus

1 L x 0.01 M = V2 x 12.39

V2 = 8.07 L

V2 = 0.807 mL

so you need to take 0.0801 mL of concentrated HCl with graduated pipet and then put it into 1 L

volumetric flask and finaly add distilled water until it reach 1 L. You can use other HCl with

another molarity that prepared in your laboratory, for the example if there is 6M HCl, use thisHCl and then calculated the solution needs using the same formula above.


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CHAPTER IV: TITRATION

4.1: INTRODUCTI ON

Titration is an analytical method that can be used to determine the amount of an unknownpresent in a sample. Titration is the process of adding a reagent solution of known concentration(the titrant) to an unknown until all of the substance of interest is used up. In order for a titrationto be useful for analysis, a number of different conditions need to be met. Among these are: areaction with known ratios of reactants and products; stable reagents for the analysis; and amethod for determining when the reaction is complete.

4.2: ACI D-BASE TITRATI ON

4.2.1: TI TRATI ON OF ACETI C ACI D

Objective: Determine the concentration of acetic acid by titration with a standard sodiumhydroxide solution.Background: Review acid/base titrations in your textbook.A weak acid is a compound that partially ionizes in aqueous solution producing hydronium(H3O+) ions. The general equation for the dissociation of any weak acid can be written as:HA (aq) + H2O (l) ??A

- (aq) + H3O+ (aq) (1)

The addition of a strong base results in a neutralization reaction in which hydroxide ions (OH-)react with hydronium to produce water:H3O

+ (aq) + OH- (aq) ? ?2 H2O (l) (2)As hydronium is consumed in the neutralization reaction, the equilibrium in equation 1 is shifted

to the right according to Le Chateliers Principle. The overall reaction for the neutralizationprocess can be written as the sum of equations (1) and (2):HA (aq) + OH- (aq) ??A- (aq) + H2O (l) (3)The concentration of the unknown solution can be determined by measuring the volume of titrantadded to reach the equivalence point. The equivalence point occurs when all of the acid has beenneutralized by the base. At the equivalence point:MAVA = MBVB (4)Where MA is the molarity of the acid, MB is the molarity of the base and VA and VB are thevolumes of the acid and base respectively. The equivalence point for the titration will bedetermined by using an indicator that changes color at the equivalence point.Phenolphthalein is pink in bases and clear in acids. When the titration mixture changes from

clear to pink, the equivalence point has been reached.Procedure:Safety goggles must be worn at all times in the laboratory.

1. Obtain about 35 mL of the acetic acid solution in a clean dry Erlenmeyer flask. Alsoobtain about 70 mL of the sodium hydroxide solution in another clean dry Erlenmeyerflask (be sure to record the concentration).

2. Rinse a 50 mL buret with two separate 5 mL portions of sodium hydroxide.3. Fill the buret with the NaOH, noting the starting volume.4. Rinse a 25 mL pipette with a small portion of the acid solution.


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Procedure:

1. Grind a tablet into a fine powder by using a mortar and pestle.2. Tare a piece of weighing paper on the balance. Carefully transfer as much powdered

sample to a piece of paper and then determine the mass.3. Place the powdered sample in a 150mL beaker.4. Add a 10.0 mL portion of ethyl alcohol to the beaker and stir.5. Add 25.0mL of water to the beaker.6. Put 3 drops of the phenolphthalein indicator in your flask. Put a magnetic stir bar in

your flask and place the flask on the center of the stir plate.7. The burette is filled with 0.100M NaOH. Make sure there are no bubbles apparent in

the burette. Record the initial volume on the burette.8. Begin titrating, Add the NaOH in 1.0mL increments, making note of when the color

change occurs. Continue adding base 5.0 mL past the equivalence point(the equivalence is approximately when the solution turned pink from the phenolphthalein).

9. Repeat steps 1-8 for the remaining tablets.

Amount of Active Ingredient in Product Tested

v Calculate the moles of base used to neutralize the acid for each aspirin.

v Acetyl salicylic acid (C9H8O4) is not a strong acid, which means that for every mole thatdissolves, not an entire mole of H+ dissociates from the acid. Nevertheless, whathydrogen ions that did dissociate were completely neutralized by the hydroxide added

from the base. How many moles of H+ were neutralized?

v For simplicitys sake, we are going to assume that acetyl salicylic acid is a strong acid,and, therefore, the initial moles of H+ equals the initial moles of acid. Since we arecomparing aspirin to aspirin, we will be able to obtain a relative comparison of theamount of acid in each aspirin. Calculate the mass of the acid for each aspirin based onthe number of moles that reacted with base.

v Check the label on the bottles and determine if your calculation in #3 is valid.

4.2.3: TI TRATION OF VINEGAR

Chemists and laboratory technicians are often called upon to make analyses of the constituents offood products. Food and drug manufacturers are required to make public the ingredients in theirproducts and often the percent of some or all of the components in a product.

Vinegar is used in many food products. Sauces, relishes, and pickles all contain large quantitiesof vinegar. Vinegar is used in food preparation as a source of acidity (tartness) and flavor.Vinegar is produced by bacterial action upon fermented juices, wine or beer. The ethanol(CH3CH2OH) in the juice or wine is oxidized to acetic acid (CH 3COOH) and minor amounts ofother organic compounds which give the characteristic flavor of each type of vinegar.

Acetic acid is the chemical in vinegar, which makes vinegar an important food additive. Thepickling industry requires vinegar of a certain percent acetic acid in order to assure pickles of


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good quality. Vinegar with a higher or lower concentration of acetic acid will make worthlesspickles. As purchased from the store, most vinegar contains 4-5% acetic acid by weight. Thismeans that in a 100g (approximately 100mL) sample of vinegar, there is 4-5g of acetic acid.

Laboratory analysis of vinegar is done by a technique known as acid-base titration. In thistechnique, the acid (here acetic acid) in the solution is reacted with a base (here sodiumhydroxide) to give a neutral salt and water.

When exactly all of the acetic acid is reacted or consumed by the sodium hydroxide, the solutionchanges from acid to neutral. We can observe this change by using an indicator which willchange color at our desired pH. If we know how much sodium hydroxide we added to neutralizethe acetic acid, then we will know the amount of acetic acid in our sample. This is because onemole of sodium hydroxide reacts with one mole of acetic acid.

We use an indicator to tell us when the titration is complete. Indicators are organic compoundsthat change color when there is a sudden change in the pH of the solution. The end point of the

titration is when a sudden change in the pH of the solution occurs. Therefore, we can tell thecompletion of the titration when we observe a change in the color of our solution to which a fewdrops of indicator have been added. The best way to monitor such a change in the pH is to usethe indicator phenolphthalein which changes color from colorless to pink at pH 8.0-9.0.

By performing a titration we can calculate the concentration of acetic acid in the vinegar. To doso we must know the volume of the acetic acid (5.00mL), the molarity of the base (the molaritywill be approximately 0.200M, record the exact molarity from the label), and the volume of thebase used to reach the end point of the titration. The volume of base will be read from the buretwhich is filled with the 0.200M sodium hydroxide. At the beginning of the titration the buret isfilled to its maximum capacity. The base is then added slowly (dropwise) from the buret to the

vinegar in an Erlenmeyer flask. Continuous swirling ensures proper mixing. The titration isstopped when the indicator shows a permanent pink coloration. The buret is read again. Thevolume of the base added is the difference between the initial volume (50.00mL) and the volumeleft in the buret at the end of the titration.

Equipment/Materials:

pH meter buffer pH 7 & 10NaOH solution vinegardroppers stirrer and stir bar buret buret clamp

NaOH +

O

OH

O

ONa

+ H 2O

sodium hydroxide sodium acetate water acetic acid


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beaker (100, 250 mL) pipet 10mLfunnel analytical balancegraph paper

Procedure:

There will be two brands of vinegar in the lab. Work in pairs. Each member of the pair willtitrate one of the two vinegar samples

1. Rinse a 50mL burette with about 5mL of the 0.200M sodium hydroxide solution. Discard therinse solution in the labeled waste container. After rinsing, fill the burette with the 0.200Msodium hydroxide solution. Use a clean dry funnel to fill the burette most of the way. Tilt theburette at a 45 degree angle, and slowly turn the stopcock to allow the solution to fill the tip

(perform this operation over a beaker). If you do not succeed the first time, repeat the operationuntil the liquid in the burette forms a continuous column from top to bottom. Clamp the buretteonto a ring stand using a burette clamp. Bring the memiscus exactly to the 50.00mL mark usinga dropper. Record theexact concentration of thesodium hydroxideon your report.

2. Use a 5mL volumetric pipette and pipette filler to add 5.00mL of vinegar to a 125mLErlenmeyer flask. Allow the vinegar to drain completely from the pipette by holding so that itstip touches the wall of the flask. Record thevolumeof vinegar and theinitial burettereadingon your report.

3. Add 3 to 5 drops of phenolphthalein to the flask with the vinegar. Also add about 10 mL ofdeionized water to the flask using your graduated cylinder to measure the water. The water isadded to dilute the natural color that some commercial vinegars have. In this way, the naturalcolor of the vinegar will not interfere with the color change of the indicator.

4. While holding the neck of the Erlenmeyer flask in your left hand, and swirling it, open thestopcock of the burette slightly with your right hand and allow the dropwise addition of the baseto the flask.

5. Using your pH meter, record pH continuously per 1 mL of sodium hydroxide added.

6. At the point where the base hits the vinegar solution the color may turn temporarily pink butthis color will disappear upon mixing the solution by swirling. Continue the titration until a faint

permanent pink color appears. Stop the titration. Record the level of the sodiumhydroxide intheburetteon your report. Be careful not to add too much base, an error called "over titration".If the indicator in your flask turns deep pink or purple, you have over titrated and will need torepeat the entire titration with a new sample of vinegar.

7. Repeat this procedure (steps 1-4) two more times with fresh 5.00 mL sample of vinegar eachtime. Record your data fromthis second run onyour report sheet.

8. Calculate the molarity of the vinegar for all three titrations.


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9. Average the three molarities. Calculate the percent concentration of acetic acid in your vinegarsample. Compare your results with the vinegar composition listed on the label of the vinegaryou titrated.

10. Obtain the data for the alternate brand of vinegar (the brand that you did not titrate) fromyour lab partner.

Titration of Vinegar - Pre-Laboratory

1. Write the balanced equation for the neutralization reaction and calculate the number ofmoles of vinegar neutralized.

2. Why was the buret rinsed with NaOH?

3. Calculate the molar mass of acetic acid. (HC2H3O2)

4. Calculate the number of grams of acetic acid in the vinegar

5. What is the Molarity of an acid, if it takes 25.27 mL of 0.198 M NaOH to titrate 5.00 mLof the acid.

6. Calculate the percent acetic acid in the vinegar from question 2

7. Graph the pH (y) vs volume NaOH (x).


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Titration of Vinegar - Data and L aboratory Report

Name______________________________ L ab Day_____________

1st run:

buret reading initial: _____________________

buret reading final: ______________________

VB: ___________________________________

ML NaOH pH ML NaOH pH


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2nd run:



VB: ___________________________________

ML NaOH pH ML NaOH pH


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3rd run:



VB: ___________________________________

ML NaOH pH ML NaOH PH


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Brand: ______________________ Manufacturers claim for % acidity:

1st run 2nd run 3rd runMB

VB

VA

MA

%

1. What is the average % acetic acid in each sample? How do these results compare with labelvalues (compare by using the percent difference)?

2. If your acetic acid sample contained a small amount of an alkaline impurity, would youexpect the molarity of this sample to be the same, greater or less than a sample that contains thesame amount of acetic acid, but no alkaline impurity?

3. If you over titrated your vinegar sample (the color was dark pink), would the calculated

molarity of the acetic acid be the same, larger or smaller as the solution if it was not overtitrated?

4.3: COMPL EXOMETRIC TIT RATION

4.3.1: TI TRATI ON OF CAL CIUM AND MAGNESIUM I N WATER

PURPOSE

The purpose of this experiment is to determine the hardness of water by measuring theconcentrations of calcium and magnesium in water samples by complexometric titration.

Determining water hardness by EDTA

All natural waters contain dissolved cations and anions. Water dissolves many ions as it flowsthrough minerals. Although water hardness is defined as the quantity of cations with a +2 or +3charge, calcium ion and magnesium ion are the most common of such ions in natural water. Theformation of solid calcium carbonate is an endothermic process. Thus, when water containingboth carbonate and calcium ions are heated, calcium carbonate can precipitate out onto the wallsof pipes, boilers, and household items such as tea pots. This can shorten the life-time of some ofthese items.


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In addition, an insoluble scum develops when hard water comes into contact with soap. Bothcalcium and magnesium ions are responsible for this precipitate. This scum can be very difficultto clean.However, there is some evidence that hard water has beneficial health effects. Selenium, forexample, may help prevent cancer. Soft water drinking supplies have been associated with anincreased heart attack risk. The quantity of hardness ions will be determined by titration.EDTA, a weak acid, will be used as the titrant. In its ionized form, it is able to form solublecomplexes with calcium and magnesium cations. The indicator added to the sample isEriochrome Black T. Initially, the indicator will form a complex with the cations. Whencomplexed it is red in color. As the EDTA is added dropwise to the sample, it replaces the Erio Tand forms more stable complexes with calcium and magnesium. When the indicator is releasedby the metal ions, it has a distinct blue color. Therefore, the endpoint of the titration is marked bythe color change from red to blue.

Procedure

v Pipet 25-ml of the water sample into an Erlenmeyer flask and dilute to a total volume ofapproximately 50 ml. Add at least one ml of pH 10 buffer solution (1/2 of a Beral pipet)to the sample. The pH should be 10. To check pH, standardize pH meter.

v Place the magnetic stirrer in the beaker and turn on the stirrer slowly; making sure thatthe bar does not hit the electrode.

v Add a few drops Eriochrome Black T indicator to the Erlenmeyer.

v Fill the burette with EDTA solution 0.01M. Record the initial burette reading.

v Immediately begin to titrate the sample two drops at a time. Be careful to titrate slowlynear the endpoint, as the color will take about 5 seconds to develop. Thus, add the lastfew drops at 3-5 second intervals. The endpoint color is blue.

v Record the initial and final burette reading to the nearest 0.1 mL.

CALCULATIONS

v Hardness is expressed as parts per million (mg per liter) of equivalent CaCO3. Forexample, if the titration required 5 ml EDTA, the calculation would be:

v Calculate the hardness of your sample in ppm of calcium carbonate

Compare the hardness of your sample with the founding of others


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4.3.2: DETERM INATIONS OF AL UMINUM AND MAGNESIUM IONS IN DRUGTA BLET (ANTACI DS)

4.3.2.1INTRODUCTION

Antacids are useful to relieve acid indigestion, upset and sour stomach, or heartburn. Antacidsmay be divided into two classes: (a) chemical antacids work by chemical neutralization of gastricacid, for example sodium bicarbonate; and (b) adsorptive antacids act by adsorbing the acid,including aluminum and magnesium salts, and calcium carbonate. The former category showsthe most rapid action, but may cause "acid rebound," a condition in which the gastric acid returnsin greater concentration after the drug effect has stopped. The latter category is less prone to therebound effect.Antacids containing aluminum ion prescribed with a low-phosphate diet can treathyperphosphatemia or prevent formations of phosphate urinary and kidney stones. Antacids

containing magnesium hydroxide and oxide with a larger dosage than normally required canproduce a laxative effect. Antacids with aluminum and magnesium hydroxides, or aluminumhydroxide alone effectively prevent significant stress ulcer bleeding in postoperative patients orthose with severe burns. Antacids including calcium ion use as diet supplements to preventosteoporosis, but constipation and renal stone formation side effects may develop.According to the information in the Medline Plus of the National Library of Medicine (seethe More Readings) related to antacid production in the U.S., there are 80 common brands.Antacids with aluminum used as the active ingredient includes alumina, aluminum hydroxide,and basic aluminum carbonate. Antacids with magnesium ion usually contain magnesia, andmagnesium trisilicate, magnesium carbonate, magnesium alginate, magnesium hydroxide, andmagaldrate, Antacid with calcium ion ordinarily has calcium carbonate. Among these brands,

major productions consist of aluminum and magnesium ions.

4.3.2.2 ANAL Y TI CAL METHODS

In this experiment, the aluminum ion and magnesium ion contents in commercial antacid will bedetermined. First, the antacid tablets are weighted, dissolved in dilute nitric acid with heating,and then diluted with deionized water to a fixed volume. Analytical methods are designed withthree protocols based on complexometric direct and back titrations along with proper indicatorsin appropriate buffer solutions.EDTA, ethylenediaminetetraacetic acid, (in quantitative analysis usually using disodiumdihydrogen ethylenediaminetetraacetate, Na2EDTA) can form strong 1:1 complexes with most

metal ions. It is most widely used as a chelating agent in analytical chemistry.Almost all metallic ions with two valences or more may be quantitatively analyzed usingcomplexometric direct or back titration with EDTA. Complexometric back titration generallyperforms when the metallic ions form a stable complex with EDTA in a slow reaction or when ametal ion blocks an indicator. The blocked indicator cannot release metallic ions, thus no colorchange will be observable at the endpoint of complexometric direct titration. Both conditionsexist in the case of aluminum ion, thus the ion is best determined by complexometric backtitration along with heating to enhance the complexation of Al-EDTA.


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4.3.2.6REAGENTS PREPARATI ON

1. M EDTA solution: Weigh approximately 1.861 g of reagent grade disodium dihydrogenethylenediaminetetraacetate dihydrate, Na2H2EDTA2H2O (molecular weight = 372.25)to precisely 0.0001 g. Quantitatively transfer it into a 500 mL of volumetric flask, add ahalf full of deionized water, swirl to enhance its dissolution, and then dilute to thecalibration mark.Stopper the flask and mix it well by inverting and shaking repeatedly. (provided)

2. 0.01 M Standardized Zinc solution: Weigh out reagent grade zinc sulfate heptahydrate,ZnSO47H2O (molecular weight = 287.53) approximately 1.438 g to precisely 0.0001 g.Transfer it quantitatively into a 500 mL volumetric flask, add a half full of deionizedwater with swirling to dissolve it, and then dilute to the calibration mark. Stopper theflask and mix it thoroughly by inverting and shaking constantly. If a 1.4377 g of the zincsalt is used, the exact molarity of the solution should be 0.01000 M.

3. Acetate-acetic acid buffer solution: Dissolve about 20.0 g of sodium acetate trihydrate,CH3COONa 3H2O, in a 1000-mL beaker containing about 1000 mL of deionizedwater.Slowly add 6 M hydrochloric acid with stirring throughout until a pH 5.0 0.1 of thesolution is detected by a pH meter. (provided)

4. Bicarbonate-carbonate buffer solution: Dissolve about 30 g of sodium carbonate,Na2CO3, in a 1000-mL beaker containing about 1000 mL of deionized water. Slowly add6-M hydrochloric acid with stirring throughout until a pH 10.0 0.1 of the solution isdetected by a pH meter. (provided)

5. Xylenol orange indicator: Dissolve 0.10 g of the acid or sodium salt form of xylenolorange in 50 mL of absolute ethanol. The prepared indicator with lemon yellow color is

suitable for analyzing solutions at pH = 5.0. The indicator prepared from the acid form isindefinitely stable, whereas that from the salt form may use only for several months.(provided)

6. Eriochrome Black T indicator: Dissolve 0.25 g of Eriochrome Black T with 50 mL ofabsolute ethanol. The blue indicator is suitable for analyzing solutions at pH = 10.0. If theindicator appears purple in color, add dropwise a pH 10 buffer solution until the colorchanges back to blue color. (provided)

4.3.2.7 SAM PLE PREPARAT ION (PROVIDED)

1. Obtain an antacid tablet from your TA and record its brand name, active ingredient(s) and

the declared quantity of each component.2. Weigh the tablet precisely to the nearest 0.0001 g. Grind it with a clean and dried mortar

and pestle to make powder as fine as possible. Place the powder onto a weighing paperon a tare balance and precisely weigh about 0.7 g of it. Transfer the powder quantitativelyto a clean 250-mL of Erlenmeyer flask. Then add about 25-mL of demonized water and5-mL of conc. hydrochloric acid to the flask.

3. Boil the mixture for about 20 minutes on a hot plate. Place a stem funnel into the flaskmouth so that the vapor can condense quickly back to water. This simultaneously helps towash down any powder that sticks onto the flask wall. If any powder remains on the wall,wash it down with a small amount of demonized water and continue heating.


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4. Remove the flask from the hot plate and allow it to cool to near room temperature orrapidly in a water bath. Filter the mixture using gravity filtration into a 250-mLvolumetric flask.Rinse the flask and the solid on the filter paper with about 10-mL of deionized watertwice to make sure that all metallic ions are transferred into the volumetric flask.

5. Dilute the solution to the calibration mark with demonized water. Stopper the flask andmix the solution well by inverting and shaking it repeatedly. Label this solution Theantacid sample solution.

4.3.2.8EXPERIMENTAL PROCEDURE

Part A: Standardization of theEDTA Solution

1. Prepare the 0.01M standardized zinc solution according to the Reagents Preparation

Procedure.2. Pipet a 10.00 mL aliquot of 0.01 M standardized zinc solution into a 125 mL Erlenmeyer

flask followed by adding about 20 mL of the bicarbonate-carbonate buffer solution (pH10.0 0.1).

3. Add 3 drops of Eriochrome Black T indicator and mix it well. The solution should appearwine red in color.

4. Direct-titrate the solution with 0.01M EDTA solution until the color changes to pure blueat the endpoint.

5. Repeat the titration twice.6. Calculate the exact concentration of the EDTA solution.

Part B: Determination of Total Aluminumand MagnesiumContents

1. Pipet a 10.00 mL aliquot of the antacid sample solution into a 125 mL Erlenmeyer flaskfollowed by adding about 20 mL of the bicarbonate-carbonate buffer solution (pH 10.0 0.1). Transfer quantitatively a 30.00 mL aliquot of 0.01 M EDTA solution to the flaskusing a burette.

2. Boil gently the mixture for 5 min. on a hot plate to speed up the formation of Al-EDTAcomplex.Add 3 drops of Eriochrome Black T indicator and mix it well. The solution should appearpure blue in color. If the EDTA is not enough to chelate all of the metallic ionscompletely, the solution should be wine red in color at this moment. In this case, put an

additional 5.00 mL or more aliquot of the EDTA solution to this wine red solution. Boilagain until the color changesto pure blue.

3. Back-titrate the solution with standardized zinc solution until the color changes to purple-blue at the endpoint (no wine red color should persist).

4. Repeat the titration twice.5. Calculate the total millimoles of aluminum and magnesium ions in the antacid sample

solution and in the tablet.


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Part C: Determination of AloneAluminumContent

1. Pipet a 10.00 mL aliquot of the antacid sample solution to a 125 mL of Erlenmeyer flask.Add about 20 mL of the acetate-acetic acid buffer solution (pH 5.0 0.1) to mask theformation of Mg-EDTA complex. Transfer quantitatively a 20.00 mL aliquot of 0.01 MEDTA solution to the lask using a burette.

2. Boil it gently on a hot plate for 5 min. to speed up the formation of Al-EDTA complex.Add 10 drops of xylenol orange indicator and mix well. The solution should appearlemon yellow in color at this moment. If the EDTA is not enough to chelate aluminumion, the solution should be deep red color. In this case, put an additional 5.00 mL or morealiquot of the EDTA solution to this deep red solution. Boil again until the color changesto lemon yellow.

3. Back-titrate the solution with a standardized zinc solution until the color changes to lightorangeyellow at the endpoint (no deep red color should remain). If the light orange-

yellow color shortly turns back to lemon yellow, continuously titrate the solution until alight orange-yellow color persists for more than 3 minutes.Note: The turning back slowly to lemon yellow color results from the complex formationof EDTA with zinc ion at a low pH solution, which is thermodynamically more stableand kinetically slower than the Zn-indicator complexation. Thus, in this protocol slowtitration will give good results.

4 Repeat the titration twice.5 Compute the millimoles and weights of aluminum ion in the antacid sample solution and

in the tablet.

4.3.2.9DATA TREATMENT

Part A: Standardization of theEDTA Solution1. Calculate the exact concentration of EDTA solution.

Part B: Determination of Total Aluminumand MagnesiumContents2. Calculate the total concentration of aluminum and magnesium of the sample.

Part C: Determination of AloneAluminumContent3. Calculate the concentration of aluminum of the sample.4. Base on part B and C, Calculate the concentration of magnesium of the sample.

4.3.2.10 QUESTIONS:1. From your calculation, compare your own results with declared concentration for the

sample you analyzed. Explain if there are significant differences.

2. If back-titration is fast in the Part C for the determination of aluminum content alone,what would be the effect on the reading of the standardized zinc solution used and thecalculated concentration of aluminum content in the sample? Explain.What would be the effect on the calculated active ingredient values in an antacid for eachg? Explain: (i) An insufficient time of boiling the solution in the Part B. (ii) Addingindicator prior to heating in the Part C


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4.3.2.10 HAZARDS

Disodium dihydrogen ethylenediaminetetraacetate dihydrate, zinc sulfate heptahydrate,potassium acetate, and sodium carbonate may cause irritation to skin, eyes and the respiratorytract.Nitirc acid is corrosive and damageable to the skin, eyes, and respiratory tract. Triethanolaminecauses skin irritation and severe eye irritation. Ethanol is highly flammable liquid; there shouldbe no open flames in the laboratory. Avoid breathing the vapors. Use with adequate ventilation.Xylenol orange, Eriochrome Black T and calmagite may cause irritation. Avoid contact with theeyes and skin.Magnesium hydroxide, aluminum hydroxide and hydrotalcite are the active ingredients incommercial antacids, thus there is no hazard in normal handling.Solutions will spatter if heated too strongly. To avoid this, heat solutions slowly and allow onlygentle boiling. Placing a funnel into the flask mouth will prevent spatter. Immediately flush the

affected skin with plenty of water for at least 15 minutes. Protective gloves and goggles must beworn at all times. .

4.4: REDOX TI TRATI ON4.4.1: MANGANIMETRI C TITRATION

INTRODUCTION

Iron ablets are prescribed for anaemia. The iron in commercial iron tablets is in the form of Fe2+.This can be oxidised to Fe3+ by the manganate (VII) (permanganate) ion. This is the reactionwhich will form the basis of our titration.

MnO4- + 8 H+ + 5 e-? Mn2+ + 4 H2O

Fe2+? Fe3+ + e-5 Fe2+ + MnO4

- + 8 H+ ? 5 Fe3+ + Mn2+ + 4 H2Ov The purple MnO4

- ion becomes colourless when it reacts with the Fe2+ ions. The endpointof the titration (that is, when all the Fe2+ ions have been used up) is that point when thepurple colour of MnO4

- no longer disappears after addition to the iron solution.v The iron tablet also contains chalk powder, sucrose, and other minor ingredients. In order

to prepare this tablet for analysis you will need to grind it and transfer it to a volumetricflask. I recommend using 2 tablets and in addition, the tablets should be dissolved inH2SO4.

APPARATUS REAGENTSbeakers, 250 cm3 potassium manganate (VII) solutionweighing bottles sulphuric acidburettes, 50 cm3 iron tabletsmeasuring cylinder, 25 cm3 distilled waterpipettes, 25 cm3

thermometervolumetric flasks, 250 cm3

Bunsen burner, tripod & gauzeConical flasks, 250 cm3


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PROCEDURE.v Accurately note the mass of one iron tablet.v Crush these tablets to a fine powder in a mortar with a pestle.v Rinse all of the crushed tablets into a conical flask and add about 25 cm3 of dilute

sulphuric acid (1 M H2SO4 ). Add more water (the final solution should have a finalvolume of 100ml) and / or warm if you experience difficulty in getting the tablet todissolve completely.

v Titrate 10.00 mL samples of your tablets solution with 0.005 M potassium permanganate.Remember that you have reached the endpoint of the titration when the purple colour ofthe permanganate just begins to stay - this indicates that all of the Fe2+ ions have reacted.Do at least three determinations although you may have to do more to ensure that yourvalues are within 0.3 mL of each other.

v Repeat the titration accurately in the chosen manner.

QUESTIONS1. Calculate the mass of iron in one tablet.2. Calculate the percentage mass of iron in a tablet3. Compare your values with other members of the group and find average values.

The following steps will help you in your calculations

1. Determine the number of moles of MnO4- used to react with 10.0 mL Fe2+ solution

(you know the volume and concentration).

2. . Calculate the number of moles of Fe

2+

that reacted (use the reaction equation)3. . Calculate the number of moles that would have reacted in 50.0 mL of Fe2+ solution(since this contains 1 tablet).

4. The number of moles of Fe2+ is the same as the number of moles of FeSO4. Calculatethe mass of FeSO4 in 1 tablet.

4.4.2: IODOMETRY

Iodometry is a method of volumetric chemical analysis, a titration where the appearance ordisappearance of elementary iodine indicates the end point. It involves the following reaction:I2 (aq) + 2e

- 2I- (aq)

1. Titration with reducing agent

Iodine, I2, is sparingly soluble in water but dissolves in potassium iodide solution, KI (aq), inwhich it forms a complex ion, KI3, potassium iodide. This solution of Iodine and potassiumiodide is known as L ugol's solution orL ugol's iodine. First made in 1829, it was named afterthe French physician J.G.A. Lugol.

I2 (aq) + 2I- I3

- (aq)


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PROCEDUREv Put in burette an aqueous solution of sodium thiosulphate Na2S2O3 whose the

concentration is 0.1Mv Put in conical flask 20ml of aqueous solution of lugol (2g I2+6gKI)v Titrate the iodine solution by Na2S2O3 until the solution becomes pale yellowv Add some drops of starch ( the solution becomes blue)v Continue the titration until the blue color despairs.

QUESTIONSWrite down the reaction for this titration if the redox potential of involved couples are:I2/I

-, E0=0,54VS2O4

2-/S2O32- E0=0.09V

2. TIT RATION WITH OXIDIZING AGENT

In acidic solutions, Hydrogen peroxide (H2O2) reacts as an oxidizing agent and oxidizes Iodide I-

to I2. The elementary iodine formed is then reduced to iodide I- by thiosulfate. Knowing the

quantity of thiosulfate used in the titration, we can then deduce both the quantity of Iodine andthe quantity of Hydrogen peroxide in the solution.We have to consider 2 Half-reactions:1. The reduction of Hydrogen peroxide (H2O2) to H2O; Standard potential: E= 1.77V2. Oxidation of Iodide I- to Iodine I2; Standard potential: E= 0.53VThis reaction is speeded up by the addition of a catalyst known as Ammonium molybdate.By the end, we have to consider the reaction between Iodine and thiosulfate.Procedurev Transfer 10ml of aqueous solution of hydrogen peroxide by mean of pipette, 10ml of 5%

solution of potassium iodide,15ml of sulfuric acid 1M , 10ml of aqueous solution ofammonium moybdate o.1%.

v And titrate with sodium thiosulphate solution of concentration of 0.1M.v Add 3 more drops and write the final volume of Na2S2O3 used.

Questionsa) Write down the reactions involved in this titrationb) From these equations deduce the concentration of hydrogen peroxide.

Preparation of starch solutionv Weight about 5mg of soluble starch and knead carefully with cold waterv Transfer the found dough into 100ml of boiling water

v Boil the solution about 2minutes ( until the solution becomes transparent)v Give the time to starch to settle at the bottom of recipientv Use the liquid part as indicator

4.5: PRECIPITAT ION TI TRATI ONS

A precipitation titration involves the formation of precipitates during the course of a titration.The titrant reacts with the analyte forming an insoluble material and the titration continues tillthe very last amount of analyte is consumed. The first drop of titrant in excess will react with anindicator resulting in a color change and announcing the termination of the titration.


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IV.1ARGENTOMETRIC TITRATIONS

The most widely applicable precipitation titrations involve the use of silver nitrate withchlorides, bromides, iodides, and thiocyanate. Since silver is always there, precipitationtitrations are referred to as Argentometric titrations.

A. MOHR METHOD

This method utilizes chromate as an indicator. Chromate forms a precipitate with Ag+ butthis precipitate has a greater solubility (Ksp=1.1.10-12) than that of AgCl (Ksp=10-10), forexample. Therefore, AgCl is formed first and after all Cl- is consumed, the first drop ofAg+ in excess will react with the chromate indicator giving a reddish precipitate.- Before the EP, Ag+ forms a precipitate with Cl- :

Ag+ + Cl- AgCl

-At the EP, Chromate forms a precipitate with Ag+:2 Ag+ + CrO4

2- Ag2CrO4

In this method, neutral medium should be used since,

In alkaline solutions (pH>10.5), silver will react with the hydroxide ions:2 Ag+ 2 HO- Ag2O +H2OIn acidic solutions (pH< 6.5), chromate will be converted to dichromate:

Ag2CrO4 + 2 H+ Cr2O7

2- + Ag+ + H2O

Blank solution preparationv In the beaker of 250ml ,put about 40ml of distilled water and 10 drops of K2CrO4 5?v Add the solution of Ag NO3(0.05M) drop by drop until the solution becomes

brownish.v Keep the solution to locate easily the EP during titration.

Chloride titrationv By means of pipette, introduce 10ml of sodium chloride in a beaker of 250ml and

add 10drop of indicator(K2CrO4 5?|)v Titrate by a standard solution of AgNO3(0.05M). During titration agitate continuously

to have a homogenous solution.v Determine the concentration of NaCl.

This procedure is valid for the titration of all halogen

B. VOLHARD METHOD

The reactant used in this method is SCN- in form of KSCN or NH4SCN. This method utilizesFe3+ (iron III) as an indicator. Fe3+ is in form of FeCl3 or NH4Fe(SO4) 2.12 H2O.(Ferricammonium sulfate).

Direct titration


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Titration of Ag+

Before the EP, Ag+ forms a precipitate with SCN- :

Ag+ + SCN- AgSCN

-At the EP, a red color due to the addition of a drop of SCN- is obtained. The followingequation describes what happens at EP.Fe3++ 6SCN- [Fe (SCN) 6]

3-

Indirect titration or Back Titration

Back titration is an analytical chemistry technique which allows finding the unknownconcentration of a reactant by reacting it with an excess volume of another reactant of knownconcentration. The resulting mixture is then titrated back, taking into account the molarity of

the excess which was added.

For this specific demonstration, the concentration of an analyte (chloride for example) isdetermined by reacting it with a known number of moles of excess reagent (Ag NO3 silvernitrate). The excess reagent is then titrated with SCN- .The indicator used is Fe3+. At the EP,a red color resulting in the formation of [Fe (SCN) 6]

3- is obtained.PROCEDUREv Take 25 ml of analyte solution and put them in a beakerv Add 2ml of concentrated HNO3 and a known volume of Ag NO3 0.1M( this volume

must be in excess and during this practice 40ml will be used )v Add again 1ml of NH4Fe(SO4)2 12H2O and agitate about one minute to coagulate the

precipitatev Titrate the excess of Ag+ in the mixture with KSCN 0.1M until the red color appears

(this color can be observed easily when the precipitate has been deposited)v Determine the concentration of halide by taking into consideration the excess of

AgNO3 .

4.6: TI TRAGE POTENTI OMETRIQUE

INTRODUCTION

Les mthodes potentiomtriques en chimie analytiques sont bases sur la relation entre le

potentiel dune pile lectrochimique et les concentrations des espces chimiques prsentes ensolution.Lapplication de la potentiomtrie met en jeu une pile lectrochimique compose dune lectrodede rfrence qui maintient un potentiel constant et dune lectrode indicatrice qui rpond lacomposition de lchantillon. Un pont salin vite le mlange de la solution de rfrence aveccelle de lchantillon.Pendant le titrage, la solution standard (solution titrante) est ajoute par fraction ; on mesure eton enregistre le potentiel de la pile aprs chaque addition de ractif.Au tour du point dquivalence qui est repr par le changement important de E, le ractif titrantest ajout en trs petite quantit.


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NB : La mthode de dosage potentiomtrique sadapte toute sorte de dosage de ractionchimique. Elle est indique pour les solutions colore ou opaque qui ne permettent paslutilisation des indicateurs ordinaires

ETALONNAGE DE LELECTRODE DE VERRE

Le pH mtre (lectrode de verre) est une lectrode membrane slective spcifique aux ionsH3O

+ .A force ionique fixe la ddp est reli au pH et donc la concentration des ions H3O+

par la relation E= E0 -2.3RT/F pHCette formule est une variation linaire de E en fonction du pH avec la pente gale au coefficientde Nernst 2.3RT/F de 58mV 200C. Ltalonnage analytique avec des solutions de pH connu(normalement des solutions tampons) correspond bien ce comportement dit nernstien c..d.idal.La reprsentation graphique E= f( pH) est labore lors de ltalonnage analytique et la droite

obtenue est appele la droitedtalonnage.Etant donn que la sensibilit de llectrode peut varier au cours du temps (sur quelques jours),vous devez talonner votre pH-mtre . Plus vous talonnez rgulirement, plus vous obtenez desmesures prcises en accord avec lquation de NERNST (rponse nernstienne)ProcdureChaque marque de pH-mtre son protocole dtalonnage.Mais il y a un principe presque commun Principe de ltalonnage deux pointsv On rgle les potentiomtres lorsque llectrode est plonge dans une solution talon de

pH = 7 (milieu de lchelle des pH dans leau).v On rgle les potentiomtres lorsque llectrode est plonge dans une solution talon de

pH = 4

NB : Si la manipulation couvre la fois le milieu acide, neutre et basique on emploi la fois lestampons 4,7, et 10MODE DEMPLOI DE LELECTRODE DE VERREv Avant de se servir de llectrode : sortir llectrode de la solution de KCl. Rincer

dlicatement llectrode avec de leau distille. Absorber dlicatement la goutte deau quiruisselle avec un papier absorbant.

v Plonger llectrode dans la solution tudier. Veiller ce que le barreau magntiquetourne bien au dessous de llectrode et quil ne risque pas de heurter celle-ci.

v Aprs lexprience : sortir llectrode de la solution tudie. Rincer la dlicatement avecde leau distille. Absorber dlicatement la goutte deau qui ruisselle avec un papierabsorbant. Plonger llectrode dans la solution de conservation si on doit rutiliser

llectrode.v En fin de TP : rincer llectrode avec de leau distille. Absorber dlicatement la goutte

deau. Remettre llectrode dans la solution de KCl concentre. Veillez ce que la bulbede llectrode plonge bien dans cette solution.

QUESTIONS1. Tracer le graphe E=f (pH)2. Calculer la pente de ce graphe3. Pourquoi est-il ncessaire dtalonner un appareil comme lectrode de verre avant son

utilisation.


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DOSAGE POTENTIOMETRIQUE DUN POLY ACI DE

Lacide phosphorique est un triacide faible dont les pKa dans leau des trois couples acido-basiques valent respectivement 25oc :- couple acide phosphorique/anion dihydrognophosphate : H3PO4 /H2PO4

- :2.15- Couple anion dihydrogenophosphate /anion (mono)hydrogenophosphate :H2PO4

-/HPO42-:7.2

- couple anion (mono)hydrogenophosphate/anion phosphate: HPO42-/PO4

3-:12.10Les valeurs de pKa1 et de pKa2de lacide orthophosphorique sont calculs en se servant desvaleurs de pH au dpart, au premier et au deuxime point quivalent.pH au dpart : pH= (Eo-E)F /2.3RT ou pH=1/2pKa1-1/2log CaAu premier point quivalent : pH=1/2pKa1+1/2pKa2Au deuxime quivalent : pH : pH=1/2 pKa2+1/2pKa3Mode opratoire- Etalonner llectrode de verre utiliser au moyen des solutions tampons-Transvaser 30ml de lacide orthophosphate 0.05 M dans bcher de 150 ml, puis ajouter quelquesgouttes de mthylorange (1erpoint quivalent)- Titrer au moyen dune solution de NaOH 0.1M- Aprs avoir atteint le premier point quivalent, ajouter quelques gouttes de phnolphtaline- Continuer le titrage en utilisant la mme solution de NaOH jusquau virage de la solution (2me

point quivalent)- Rpter lexprience trois foisNB : Ajouter chaque fois 1ml et noter le potentiel obtenu

Dpasser le 2mepoint quivalent de 10ml au moinsQuestions

1. Montrer clairement comment vous avez procd pour prparer les diffrentes solutionsutilises.2. Donner les concentrations des espces prsentes en solution au premier et au deuxime pointquivalent.3. Tracer la courbe de titrage (la courbe drive).4. Calculer les valeurs de pKa1 et pKa2 de lacide orthophosphorique.

CHAPTER V: EL ECTROCHEMISTRY

ELECTROCHEMICAL CELL

Introduction

The device used to study reactions electrically is called an electrochemical cell. Such a cellconsists of two electrodes, or metallic conductors, dipping into an electrolyte, an ionic conductor,which may be a solution, a liquid, or a solid. An electrode and its electrolyte comprise anelectrode compartment. If the two electrolytes are different, then two compartments may bejoined by a salt bridge, which is an electrolyte solution that completes the electrical circuit bypermitting ions to move between the compartments and so enables the cell to function (see figurebellow).


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Reaction at electrodesIn electrochemical cell, the oxidation half-reaction takes place in one compartment and thereduction half reaction takes place in the other compartment. As the reaction proceeds, theelectrons released in oxidation half reaction:In one compartment travel through the external circuit, and enter the cell through the otherelectrode, where they bring about the reduction:Electrical neutrality is preserved in the electrolytes by the flow of cations and anions in theopposite directions through the salt bridge as shown in the figure above.

MATERIAL S REAGENTS

2 beakers of 500ml CuSO4Blade or rod of copper ZnSO4Blade or rod of zinc Agar-agar Digital voltmeter with high internal resistance KNO3Two external conductors in copper Distilled water U-shape tube

SALT BRIDGE PREPARATION

v In the conical flask heat 100ml of a solution obtained by mixing 30g KNO 3 ,3g of agar-agar and distilled water. This mixture should be covered and gently heated, with stirring,

_+

V

OxidationReduction

laboratory practices in progress

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